Video: Acidity and Basicity

In this video, we will define acids and bases, learn about their properties, and identify some common acids and bases we’ll encounter both in a chemistry lab and in everyday life.

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Video Transcript

In this video, we will define acids and bases, learn about their properties, and identify some common acids and bases we’ll encounter both in chemical reactions and in everyday life.

The terms acid and base might bring some scary imagery to mind, but acids and bases are all around us. Acids are well known for being corrosive to the skin, but they also have a sour taste. For instance, citric acid is what gives lemons and other citrus fruits their characteristic sour flavor. And acids in general are pretty commonly found in foods. Apples contain malic acid. Milk contains lactic acid. Vinegars all contain acetic acid. And the gastric juices inside our stomachs are made of HCl or hydrochloric acid.

Bases, on the other hand, are often bitter and have a slippery, soapy feel to them, which is why you’ll often see bases in soaps, detergents, and other cleaning products. For instance, ammonia is a common household cleaner. Sodium hydroxide, which is otherwise known as lye, is used in soap-making. It’s an oven and drain cleaner. And it’s traditionally used in some kinds of food preparations, like making pretzels and bagels. Sodium hydrogen carbonate, otherwise known as baking soda, is used in baking as a leavening agent. Sodium hydrogen carbonate can also be used as an antacid. But if you buy antacids from the store, you’re most likely buying calcium carbonate.

Sometimes you’ll see bases referred to as alkalis. But there’s a slight distinction between the two terms. An alkali is a base that can be dissolved in water. So all alkalis are bases, but not all bases are alkalis. But what makes an acid an acid and a base a base?

These examples we’ve seen are all pretty diverse in their chemical formulas. Though people and scientists have been using acids and bases for hundreds and hundreds of years, acids and bases weren’t well defined until the Swedish chemist Svante Arrhenius came up with a way to define acids and bases in 1884, which is now referred to as the Arrhenius definition of acids and bases.

He defined an acid as a substance that has hydrogen in its chemical formula and ionizes in water to form H+ ions. For example, HCl is a classic Arrhenius acid. It ionizes in water to produce H+ ions and Cl− ions. Well, there’s actually a slight complication with how this chemical formula is written since H+ ions don’t typically exist in water on their own. When H+ is in water, it will immediately react with the water to form the ion hydronium or H3O+. So H+ ions aren’t really produced when an acid ionizes. So it would actually be more correct to write HCl plus H2O reacting to form H3O+ plus Cl−. But the other way is shorter and easier, so you’ll often see it written in that way as well. But technically, acids produce H3O+ ions in water, not H+ ions.

The Arrhenius definition of a base is a substance that has OH in its chemical formula that ionizes in water to form OH− or hydroxide ions. Sodium hydroxide is a great example of an Arrhenius base. In water, it will ionize to form sodium ions and hydroxide ions.

Now, this definition isn’t just some abstract thing we have to imagine occurring in a solution that’s based off a chemical equation that we see written in a book. We can actually measure the concentration of H3O+ ions in a solution to determine whether it’s acidic or basic. When we measure the concentration of hydronium ions in the solution, we can assign that solution a pH value. A pH value of less than seven corresponds to a solution that’s acidic. And a pH value of greater than seven is a solution that’s basic. If the pH value is seven, that solution is neutral.

The Arrhenius definition of acids and bases gives us a way to categorize chemical substances as acids or bases, but it does have a couple of problems. The primary issue with this definition is that there’s plenty of chemical substances that act as bases, meaning they produce hydroxide ions in water, but they don’t have OH in their chemical formula. This definition is also reliant on the acid or base being dissolved in water. But acid–base reactions can happen between gaseous substances and in nonaqueous conditions.

To address these issues, Johannes Nicolaus Brønsted and Thomas Martin Lowry came up with a different definition in 1923. The Brønsted–Lowry definition of an acid is a chemical species that donates a proton and a base is a chemical species that accepts a proton. When we say proton here, we’re referring to an H+ ion, since hydrogen has one proton and one electron, so an H+ ion is just one proton. So according to this definition, an acid–base reaction is just a proton transfer process because one chemical species gives up a proton and the other chemical species takes it.

To get a feel for this definition, let’s take a close look at some reactions. We’ve already seen how HCl is an Arrhenius acid. But let’s take a look at how it’s also a Brønsted–Lowry acid. When HCl reacts with H2O, the proton on HCl gets donated to H2O. Since water gains this proton, it forms H3O+. And since HCl lost a proton, we’re left with Cl−.

Now, let’s take a look at a reaction with a base. In this reaction, we have ammonia reacting with water. This was the chemical species that gave us trouble with the Arrhenius definition of bases. In this reaction, ammonia, or NH3, has a lone pair, which is the perfect site to accept a proton, which it accepts from water. Once ammonia accepts this proton, it is now NH4+. And water having lost a proton is now OH−.

In these kind of acid–base reactions, we’ll always end up with chemical species that differ from each other by a proton. And we call these chemical species an acid–base conjugate pair. An acid always loses a proton to form its conjugate base. So Cl− is the conjugate base of HCl. And a base always accepts a proton to form its conjugate acid. So NH4+ is the conjugate acid of NH3.

There’s one other thing to notice here. We have water participating in both reactions. In the first, it accepts a proton from HCl, which means that according to the Brønsted–Lowry definition of acids and bases, water is acting as a base in this reaction. And in the bottom reaction with ammonia, water is losing a proton. So, according to the Brønsted–Lowry definition of acids and bases, here it’s acting as an acid.

There are many chemical species that can act as both an acid and a base. When a substance can act as both an acid and a base, it’s called amphoteric. But of the amphoteric chemical species, water is a bit unique. Since it’s both an acid and a base, it constantly reacts with itself in a process that’s known as self-ionization or autoionization. When this happens, one water molecule donates its proton to the other water molecule. The water molecule that takes the proton, which would make it a Brønsted–Lowry base, forms H3O+. The other water molecule here is acting as an acid since it donates its proton to the other water molecule, leaving it with one less proton, forming OH−. This is why pure water is neither acidic or basic, since it’s forming equal amounts of H3O+ and OH− ions. However, it does this in very small amounts. The concentration of both OH− and H3O+ will only be 10 to the minus seven molar.

So now that we know how to define acids and bases, let’s explore some of the different kinds of acids and bases we’ll encounter. Now, there are two main categories of acids and bases, strong acids or bases and weak acids or bases. This distinction essentially has to do with how many units of the acid or base ionize to form H3O+ or OH− ions.

Let’s take a look at two pretty similar-looking acids, HI and HF. Both contain hydrogen and a halogen, that is, an element in group 17. Since both HI and HF are acids, when they’re introduced to water, they will donate their protons to water, forming H3O+ and an anion, I− in the case of HI and F− in the case of HF. But if we compare the resulting anions from these acids reacting, the F− ion is smaller than the I− ion. But they have the same amount of negative charge, which means that the F− ion has a much greater density of negative charge than the I− does. Because F− has so much negative charge density, it’s highly likely that the F− ions will take a proton from the H3O+ ions, which will form HF again.

When the reaction can go in the reverse direction like this, we can indicate that with two half arrows. When this happens, the two reactions will eventually establish an equilibrium between the forward and reverse reactions. But that’s the subject for another video.

The overall result of this is that if we put the same amounts of HI and HF molecules in a solution, the solution containing HI will end up with more H3O+ ions than the solution of HF because the HI solution ionized completely and the HF solution still has some HF molecules hanging around. So we call HI a strong acid and HF a weak acid. So put simply, strong acids and bases will ionize completely in water, while weak acids and bases do not.

Now, it’s worth mentioning just because an acid or base is strong, that doesn’t necessarily mean we’ll have more H3O+ or OH− ions in a solution than a weak acid or base would. Concentration plays an important role here as well. We’ve just seen now if we have the same amounts of HI and HF molecules, that is, the same concentration of both, assuming they’re in the same volume of water, we’ll end up with more H3O+ ions in the solution of HI than in the solution of HF because HI is a strong acid. But what if our solutions had different concentrations? Specifically, what if the strong acid solution was way less concentrated than the weak acid solution?

In this case, even though the strong acid completely ionizes, there won’t be as many acid molecules to ionize. So there just won’t be that many H3O+ ions in the solution. In this case, the weak acid will comparatively have a higher concentration of H3O+ ions even though it’s a weak acid.

With all this in mind, let’s take a look at some common strong and weak acids and bases. Let’s start with the acids. Let’s start with the strong acids because there’s only a handful of them. Unfortunately, this is a list that needs to be memorized. Our first three strong acids are acids with halogens in them: HCl or hydrochloric acid, HBr or hydrobromic acid, and HI or hydroiodic acid. Then we have four oxoacids, that is, acids that contain oxygen: HNO3, nitric acid; H2SO4, sulfuric acid; HClO4, perchloric acid; and HClO3, or chloric acid. Chloric acid is sometimes not included as a strong acid, depending on who you ask. So that’s just something to look out for.

Then we have the weak acids. There’s many weak acids. So we’ll just cover a few of them, like hydrofluoric acid, which we’ve already seen, HCN or hydrogen cyanide, and H2S or hydrogen sulfide. Then we have some oxoacids that are weak acids, like HClO or hypochlorous acid, H3PO4 or phosphoric acid, and HNO2 or nitrous acid. Then we have all the carboxylic acids, which are molecules that contain this group that has a carbon double bonded to an oxygen and single bonded to an OH. An example of this is acetic acid, which was the acid that’s in vinegar.

As we look through these lists, we’ll notice that most of these acids only contain one hydrogen, but some of them contain more. For instance, H2SO4 contains two hydrogens and H3PO4 contains three. Acids that can only donate one proton are called monoprotic or monobasic acids. There are plenty of examples of monoprotic acids on the screen right now, including acetic acid. Even though it has more than one proton, only one of them can be donated. So it’s still a monoprotic acid. Acids that have two protons to donate are called diprotic or dibasic, like sulfuric acid and hydrogen sulfide. And acids that have three protons that can be donated, like H3PO4, are referred to as triprotic or tribasic.

Now, let’s take a look at the bases. The strong bases are much more regular than the list of strong acids. They will be compounds containing O2− or OH− plus either a group one metal — lithium, sodium, potassium, rubidium, or cesium — or some of the group two metals — calcium, strontium, and barium. So some examples of strong bases would be NaOH, Li2O, Ca(OH)2, or BaO. The weak bases are mostly nitrogen-containing compounds, like NH3 or ammonia, like we’ve seen.

Then we have a class of compounds that are referred to as amines. Examples of these would be the compounds methylamine, dimethylamine, and trimethylamine. Other examples of amines that are weak bases are the neurotransmitters in the human body, like dopamine.

Now, we’ve covered a lot of information in this video, so let’s summarize with the key points. The Arrhenius definition of acids is a substance that contains hydrogen and yields H+ ions in water, a classic example being hydrochloric acid. And the Arrhenius definition of a base is a substance that contains OH and yields OH− ions in water, a classic example being sodium hydroxide. In the Brønsted–Lowry definition of acids and bases, an acid is the proton donor. For example, hydrofluoric acid donates a proton to water, which forms the fluoride anion and hydronium. And a base is a proton acceptor. For instance, ammonia accepts a proton from water, giving us NH4+ and OH−. Water can act as both an acid and a base. And it can react with itself in a process known as autoionization or self-ionization. An acid or base is a strong acid or base if it ionizes completely in solution. But if the acid or base doesn’t completely ionize, it’s a weak acid or base.

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